#textbook #chemistry #bonding #metalicbonding

Pg:- 65,66,67,68

Ch:- 3

It is an electrostatic force of attraction between the mobile sea of electrons and the regular array of positive metal ions within a solid metal.

Physical properties of metals:

• Most metals have high melting and boiling points.

• Metals are good conductors of electricity.

• Metals are easily bent and shaped or stretched into wires.

• Delocalized Electrons ("Electron Sea Model"):

  • In metallic bonding, the outermost (valence) electrons of metal atoms are not bound to any specific atom. Instead, these electrons are free to move throughout the entire metal lattice.

  • This is often referred to as the "electron sea model," where metal atoms (positive metal cations) exist in a fixed, orderly lattice, while the delocalized electrons move freely within the lattice.

  • The delocalization of electrons leads to strong attractions between the metal cations and the electrons, holding the structure together.

• Metal Lattice Structure:

  • Metals are arranged in a crystal lattice structure, where the metal atoms are closely packed in a regular pattern. The lattice is held together by the attraction between the positively charged metal ions and the negatively charged delocalized electrons.

  • Common lattice types in metals include body-centered cubic (BCC), face-centered cubic (FCC), and hexagonal close-packed (HCP).

Factors Affecting the Strength of Metallic Bonding

1. Number of Delocalized Electrons:

   • Metals with more delocalized electrons available for bonding (i.e., those with more valence electrons) tend to form stronger metallic bonds. For example, magnesium (Mg) with two valence electrons per atom forms stronger bonds than sodium (Na) with one valence electron per atom.

2. Size of Metal Ions:

   • Smaller metal ions can pack more closely together, resulting in stronger metallic bonds. For example, lithium (Li) forms weaker bonds than magnesium (Mg) because the Li⁺ ions are larger and more spaced out in the lattice.

3. Charge of Metal Ions:

   • Metals that form ions with higher charges will have stronger metallic bonds. For example, magnesium (Mg²⁺) forms stronger bonds than sodium (Na⁺), because the doubly charged Mg²⁺ ion attracts the delocalized electrons more strongly than the singly charged Na⁺ ion.

Examples of Metallic Bonding

1. Sodium (Na):

   • Sodium has a single valence electron that is delocalized within the metal lattice. The bonding in sodium is relatively weak, giving sodium its low melting point (98°C) and soft texture.

2. Magnesium (Mg):

   • Magnesium has two valence electrons that are delocalized in the lattice. The additional electrons and smaller ion size compared to sodium lead to stronger metallic bonding. This gives magnesium a higher melting point (650°C) and a harder texture.

3. Iron (Fe):

   • Iron has more delocalized electrons and a more densely packed lattice structure, leading to very strong metallic bonds. This makes iron strong, hard, and a good conductor of heat and electricity. Its ability to form alloys like steel further enhances its mechanical properties.

Important points:

• Electron Sea Model: Metal atoms are surrounded by a sea of delocalized electrons.

• Strong Electrostatic Attraction: The attraction between positive metal ions and the free electrons creates strong metallic bonds.

• Conductivity: Metals conduct electricity and heat well due to the mobility of free electrons.

• Malleability and Ductility: Metals can be shaped and drawn into wires due to the flexibility of the metal lattice.

• High Melting and Boiling Points: Strong metallic bonds require significant energy to break.

• Luster and Reflectivity: Metals are shiny because free electrons can absorb and re-emit light.